Acids & Bases Acids & Bases

A More Detailed Look A More Detailed Look

VCE Chemistry VCE Chemistry

Unit 2: Environmental ChemistryUnit 2: Environmental Chemistry

Area of Study 1 – WaterArea of Study 1 – Water

• In early times, acids and bases were defined in terms of their properties, such as their taste, their ability to corrode metals and the change they caused in the colours of indicators.

• In the late 19th century, the Swedish chemist Svanté Arrhenius attempted to explain the reactions between acids and bases in terms of the particles they released in their aqueous solutions.

• He proposed that, in aqueous solution, a base releases the hydroxide ion, OH-, and an acid releases the hydrogen ion, H+.

• In neutralisation reactions, these ions combine to produce water:

H+ (aq) + OH- (aq) → H2O (l)

• Arrhenius’ theory was limited to reactions in aqueous solution and did not explain why acids and bases released these particles.

Arrhenius Acids & Bases Arrhenius Acids & Bases

• A more useful definition used today was first proposed independently by the Danish chemist Johannes Brønsted and the English chemist Thomas Lowry in 1923.

• Brønsted and Lowry described reactions of acids as involving the donation of a hydrogen ion (H+).

• A hydrogen ion is a hydrogen atom that has lost its only electron.

• In most cases a hydrogen ion is a proton.

• Chemists use the terms hydrogen ion and proton interchangeably.

The Lowry-Brønsted TheoryThe Lowry-Brønsted Theory

• According to the Lowry–Brønsted theory:

An acid is a substance that can donate a proton (H+).

A base is any substance that can accept protons from an acid.

• No substance can act as an acid unless there is a base present to accept its protons.

The Lowry-Brønsted TheoryThe Lowry-Brønsted Theory

• As protons are exchanged from an acid to a base, this definition explains why acids and bases react together.

• For example, hydrogen chloride (HCl) is a molecular compound that is very soluble in water.

• In an aqueous solution of hydrogen chloride (hydrochloric acid), nearly all the hydrogen chloride is present as ions (HCl is ionised).

• In this reaction, each hydrogen chloride molecule has donated a proton to a water molecule.

• According to the Brønsted–Lowry theory, the hydrogen chloride has acted as an acid.

• The water molecule has accepted a proton from the hydrogen chloride molecule, so has acted as a base.

HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq)

The Lowry-Brønsted TheoryThe Lowry-Brønsted Theory

• Some substances can behave as either acids or bases, depending what they are reacting with, and can therefore donate or receive protons.

• Such substances are said to be amphiprotic.

• For example, the hydrogen carbonate ion (in bicarbonate soda or fizzy soft drinks) can potentially react with water molecules in two possible ways:

• As an acid: HCO3- (aq) + H2O (l) → H3O+ (aq) + CO3

2- (aq)

• As a base: HCO3- (aq) + H2O (l) → H2CO3 (aq) + OH- (aq)

Amphiprotic SubstancesAmphiprotic Substances

• Experiments show that different acid solutions of the same concentration do not have the same pH.

• Some acids donate a proton more readily than others.

• The Brønsted–Lowry theory describes the strength of an acid as its ability to donate hydrogen ions to a base.

• The strength of a base is a measure of its ability to accept hydrogen ions from an acid.

A strong acid donates protons readily.

A strong base accepts protons readily.

Weak acids or bases do not donate or accept protons readily.

Acid & Base StrengthAcid & Base Strength

Acid & Base StrengthAcid & Base Strength

• Acids that ionise completely in solution are called strong acids.

• Strong acids donate protons easily, and their solutions would contain ions, with virtually no unreacted acid molecules remaining.

• Hydrochloric acid, sulphuric acid and nitric acid are examples of strong acids.

HCl (g) + H2O (l) → H3O+ (aq) + Cl− (aq)

H2SO4 (l) + H2O (l) → H3O+ (aq) + HSO4− (aq)

HNO3 (l) + H2O (l) → H3O+ (aq) + NO3− (aq)

Strong AcidsStrong Acids

• Vinegar is a solution of ethanoic acid.

• Pure ethanoic acid is a polar covalent molecular compound that ionises in water to produce hydrogen ions and ethanoate ions.

• In a 1.0 M solution of ethanoic acid, only a small proportion (less than 1%) of the ethanoic acid molecules are ionised at any one time.

• We indicate its weaker ability by writing an equation for its reaction with water with a ‘reversible’ arrow (⇋) as shown below:

CH3COOH (aq) + H2O (l) ⇋ CH3COO- (aq) + H3O+ (aq)

• Ethanoic acid is, therefore, described as a weak acid in water.

Weak AcidsWeak Acids

• The ionic compound sodium oxide (Na2O) dissociates in water, releasing sodium ions (Na+) and oxide ions (O2−).

• The oxide ions react completely with the water, accepting a proton to form hydroxide ions (OH−):

O2− (aq) + H2O (l) → OH− (aq) + OH− (aq)

• The oxide ion is an example of a strong base.

• Strong bases accept protons easily.

Strong BasesStrong Bases

• Ammonia is a covalent molecular compound that ionises in water by accepting a proton.

• This ionisation can be represented by the equation:

NH3 (aq) + H2O (l) ⇋ NH4+ (aq) + OH− (aq)

• Ammonia is behaving here as a base because it has gained a proton.

• Water has donated a proton and so is behaving as an acid.

• Only a small proportion of ammonia molecules ionise so that a 1.0 M solution of ammonia contains mostly ammonia molecules together with some ammonium ions and hydroxide ions.

Weak BasesWeak Bases

• Some acids are capable of donating more than one proton from each molecule and are said to be polyprotic.

• The number of hydrogen ions an acid can donate depends on the structure of the acid.

• Monoprotic acids can donate only one proton and include hydrochloric acid (HCl), nitric acid (HNO3) and ethanoic acid (CH3COOH).

• Diprotic acids, such as sulphuric acid (H2SO4) and carbonic acid (H2CO3), can donate two protons.

• Triprotic acids can donate three protons and include phosphoric (H3PO4) and boric (H3BO3) acid.

• Polyprotic acids do not donate all their protons at once, but do so in steps when reacting with a base:

H2SO4 (aq) + H2O (l) → HSO4- (aq) + H3O+ (aq)

HSO4- (aq) + H2O (l) ⇋ SO4

2- (aq) + H3O+ (aq)

Polyprotic AcidsPolyprotic Acids

• The strength of an acid is different from the concentration of an acid.

• The strength of a solution is determined by the number of ions present.

• A strong acid is completely ionised in solution.

• Concentration, however, refers to the amount of an acid or base that is dissolved in a given volume of water.

• A large amount will always produce a concentrated solution whereas a small amount in the same volume of water produces a dilute solution.

• It is possible to have a weak, concentrated acid or a dilute solution of a strong acid.

Strength versus ConcentrationStrength versus Concentration

• The acidity of a solution is a measure of the concentration of hydrogen ions present.

• The higher the concentration of hydrogen ions, the more acidic the solution.

• Water is a molecular compound that has the ability to act as both an acid and a base (amphiprotic).

• Pure water undergoes self-ionisation to a very small extent, but enough for pure water to conduct electricity slightly.

• This reaction can be represented by the equation:

H2O (l) + H2O (l) ⇋ H3O+ (aq) + OH− (aq)

• In this reaction, water behaves as a very weak acid and a very weak base, producing one hydrogen ion (H3O+) for every hydroxide ion (OH−).

• The concentration of these ions is very low in pure water (10-7 M for each specie at 25°C).

Acidic, Basic & Neutral SolutionsAcidic, Basic & Neutral Solutions

• Pure water is a neutral substance because the concentration of H3O+ ions is equal to the concentration of OH− ions present.

• If an acid is added to water, however, more H3O+ ions are produced.

This results in an acidic solution.

• Similarly, if a base is added to water, more OH− ions are produced.

This results in a basic solution.

• Therefore:

Acidic solutions contain a greater concentration of H3O+ than OH−

Neutral solutions contain equal concentrations of H3O+ and OH−

Basic solutions contain a lower concentration of H3O+ than OH−.

Acidic, Basic & Neutral SolutionsAcidic, Basic & Neutral Solutions

• The pH scale is a useful way of indicating the acidity of a solution.

• Mathematically, pH is defined as:

pH = −log10 [H3O+]

where [H3O+] is measured in mol/L or (M).

• If we know the pH of a solution we can calculate, its concentration using:

[H3O+] = 10-pH

• Since the pH scale is based upon the negative logarithm of the hydrogen ion concentration, the pH of a solution decreases as the concentration of hydrogen ions increases.

• Since the pH scale is a logarithmic scale, increasing the concentration of H+ by a factor of 10 results in a decrease of one pH unit.

The pH ScaleThe pH Scale

The pH ScaleThe pH Scale

The pH ScaleThe pH Scale

1. What is the pH of a solution in which [H+] = 0.0135 M?

A. pH = -log10 [H+]

pH = -log10 (0.0135)

pH = - (-1.87) = 1.87

2. What is the concentration of i) hydronium ions and ii) hydroxide ions in a solution of pH 12.3

Ai [H3O+] = 10-pH

[H3O+] = 10-12.3

[H3O+] = 5.01 x 10-13 mol/L

Aii [OH-] = 10-14 / [H3O+]

[OH-] = 10-14 / 5.01 x 10-13 = 0.02 mol/L

[OH-] = 0.02 mol/L

Calculating pH - ExamplesCalculating pH - Examples